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Chapter 10: Chemical Bonding II – Molecular Shapes, Valence Bond Theory, and Molecular Orbital Theory

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Chapter 10: Chemical Bonding II

Introduction

This chapter explores advanced concepts in chemical bonding, focusing on molecular shapes, the Valence Shell Electron Pair Repulsion (VSEPR) theory, valence bond theory, and molecular orbital theory. Understanding these models is essential for predicting molecular geometry, bond angles, and the properties of chemical substances.

Molecular Shape and Its Importance

Taste and Molecular Interaction

The interaction between food molecules and taste cells is determined by the molecule’s shape and charge distribution. Molecules must fit into the active sites of proteins on taste cells, triggering nerve signals. This principle illustrates the importance of molecular geometry in biological systems.

  • Example: Sugar and artificial sweeteners interact with T1r3 receptor proteins. Artificial sweeteners may bind more strongly, making them taste "sweeter."

Valence Shell Electron Pair Repulsion (VSEPR) Theory

Basic Principles

VSEPR theory is a model used to predict the shapes of molecules based on the repulsion between electron groups around a central atom. Electron groups include lone pairs, single, double, and triple bonds.

  • Electron groups repel each other through coulombic forces.

  • The most stable arrangement is when electron groups are as far apart as possible.

Electron Groups

The Lewis structure helps determine the number of electron groups around a central atom:

  • Each lone pair = 1 electron group

  • Each bond (single, double, triple) = 1 electron group

  • Example: In NO2, nitrogen has three electron groups: one lone pair, one single bond, one double bond.

Electron Group Geometry

There are five basic arrangements of electron groups around a central atom, leading to five basic electron geometries. Resonance does not affect the electron geometry.

Basic Electron Geometries

Electron Groups

Geometry

Bond Angles

Example

2

Linear

180°

CO2, BeCl2

3

Trigonal Planar

120°

BF3

4

Tetrahedral

109.5°

CH4

5

Trigonal Bipyramidal

120°, 90°

PCl5

6

Octahedral

90°

SF6

Linear Geometry (2 Electron Groups)

  • Electron groups on opposite sides of the central atom

  • Bond angle: 180°

  • Example: CO2, BeCl2

Trigonal Planar Geometry (3 Electron Groups)

  • Electron groups form a triangle around the central atom

  • Bond angle: 120°

  • Example: BF3

Tetrahedral Geometry (4 Electron Groups)

  • Electron groups form a tetrahedron

  • Bond angle: 109.5°

  • Example: CH4

Trigonal Bipyramidal Geometry (5 Electron Groups)

  • Electron groups form two tetrahedra base-to-base

  • Axial positions: above and below central atom

  • Equatorial positions: same plane as central atom

  • Bond angles: 120° (equatorial), 90° (axial-equatorial)

  • Example: PCl5

Octahedral Geometry (6 Electron Groups)

  • Electron groups form two square-base pyramids base-to-base

  • All positions equivalent

  • Bond angle: 90°

  • Example: SF6

The Effect of Lone Pairs

Bond Angle Distortion

Lone pairs exert greater repulsion than bonding pairs, causing bond angles to decrease from ideal values. Repulsion order: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.

Molecule

Lone Pairs

Bond Angle

CH4

0

109.5°

NH3

1

107°

H2O

2

104.5°

Derivatives of Tetrahedral Geometry

  • One lone pair: Pyramidal shape (e.g., NH3)

  • Two lone pairs: Bent shape (e.g., H2O)

Derivatives of Trigonal Bipyramidal Geometry

  • One lone pair: Seesaw shape

  • Two lone pairs: T-shaped

  • Three lone pairs: Linear shape

  • Lone pairs occupy equatorial positions for more space

Derivatives of Octahedral Geometry

  • One lone pair: Square pyramidal shape

  • Two lone pairs: Square planar shape

  • Lone pairs take positions opposite each other

Predicting Molecular Geometry

  1. Draw the Lewis structure.

  2. Determine the number of electron groups around the central atom.

  3. Classify each electron group as bonding or lone pair.

  4. Use the electron group geometry to predict molecular shape and bond angles.

Representing Three-Dimensional Shapes

  • Central atom is placed in the plane of the paper.

  • Atoms in the same plane: straight lines.

  • Atoms in front: solid wedge.

  • Atoms behind: hashed wedge.

Molecular Polarity

Dipole Moments and Polarity

Molecular polarity depends on both the polarity of individual bonds and the overall molecular geometry. A molecule is polar if it has a net dipole moment.

  • Example: HCl is polar; CO2 is nonpolar despite polar bonds due to its linear geometry; H2O is polar due to its bent shape.

Predicting Polarity

  1. Draw the Lewis structure and determine molecular geometry.

  2. Identify polar bonds.

  3. Determine if polar bonds add to give a net dipole moment.

Polarity and Solubility

  • Polar molecules dissolve well in polar solvents like water.

  • Ionic compounds also dissolve well in water.

  • Molecules with both polar and nonpolar regions may have unique solubility properties.

Problems with Lewis Theory

  • Does not provide accurate numerical predictions for bond strength and length.

  • Cannot always predict actual bond angles or magnetic properties.

  • Struggles with resonance and delocalization.

Valence Bond Theory

Basic Concepts

Valence Bond (VB) theory extends the quantum-mechanical model to chemical bonding. Bonds form when atomic orbitals overlap, and electrons are spin-paired.

  • Hybridized atomic orbitals are combinations of standard atomic orbitals.

  • The geometry of overlapping orbitals determines molecular shape.

Hybridization

  • Atoms hybridize orbitals to maximize bonding and stability.

  • Types: sp, sp2, sp3, sp3d, sp3d2

  • Number of hybrid orbitals formed equals the number of atomic orbitals combined.

  • Example: Carbon in methane (CH4) uses sp3 hybridization for four equivalent bonds at 109.5°.

Types of Bonds

  • Sigma (σ) bond: Orbitals overlap along the axis connecting nuclei. Stronger bond.

  • Pi (π) bond: Orbitals overlap side-by-side, perpendicular to the axis. Weaker than σ bonds.

Bond Rotation

  • σ bonds allow free rotation around the bond axis.

  • π bonds restrict rotation due to side-by-side overlap.

Molecular Orbital (MO) Theory

Basic Principles

MO theory applies quantum mechanics to molecules, creating molecular orbitals that are delocalized over the entire molecule. Electrons belong to the whole molecule, not just individual atoms.

  • Constructive combination: Bonding molecular orbital (lower energy)

  • Destructive combination: Antibonding molecular orbital (higher energy)

Bond Order

Bond order indicates the strength and stability of a bond:

  • Higher bond order = stronger, shorter bond

  • If bond order = 0, no bond forms

Magnetic Properties

  • Unpaired electrons in MO diagrams indicate paramagnetism (e.g., O2).

  • All electrons paired: diamagnetic

Heteronuclear Diatomic Molecules

  • Atomic orbitals of different energies contribute unequally to molecular orbitals.

  • Lower-energy atomic orbitals contribute more to bonding MOs.

  • Nonbonding MOs remain localized on the atom donating the orbital.

Polyatomic Molecules

  • All atomic orbitals combine to form molecular orbitals delocalized over the entire molecule.

  • MO theory better matches real molecular properties than Lewis or VB theory.

Summary Table: Electron Group Geometry and Hybridization

Electron Groups

Geometry

Hybridization

Bond Angles

2

Linear

sp

180°

3

Trigonal Planar

sp2

120°

4

Tetrahedral

sp3

109.5°

5

Trigonal Bipyramidal

sp3d

120°, 90°

6

Octahedral

sp3d2

90°

Additional info: These notes expand on the provided slides with definitions, examples, and tables for clarity and completeness, suitable for exam preparation in General Chemistry.

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